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Mechanisms of Sb(III) oxidation by pyriteinduced hydroxyl radicals and hydrogen peroxide Linghao Kong, Xingyun Hu, and Mengchang He Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 25 Feb 2015 Downloaded from http://pubs.acs.org on February 25, 2015
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Mechanisms of Sb(III) oxidation by pyrite-induced hydroxyl radicals
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and hydrogen peroxide
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Linghao Kong, Xingyun Hu, Mengchang He*
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State Key Laboratory of Water Environment Simulation, School of Environment, Beijing Normal
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University, Beijing 100875, P.R. China.
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* Corresponding author:
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Mengchang He
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State Key Laboratory of Water Environment Simulation
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School of Environment
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Beijing Normal University
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Beijing 100875, P.R. China.
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Tel : +86-10-5880 7172
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Fax: +86-10-5880 7172
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E-mail:
[email protected] (MC He)
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TOC ART
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ABSTRACT
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Antimony (Sb) is an element of growing interest, and its toxicity and mobility are
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strongly influenced by redox processes. Sb(III) oxidation mechanisms in pyrite
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suspensions were comprehensively investigated by kinetic measurements in oxic and
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anoxic conditions and simulated sunlight. Sb(III) was oxidized to Sb(V) in both
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solution and on pyrite surfaces in oxic conditions; the oxidation efficiency of Sb(III)
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was gradually enhanced with the increase of pH. The pyrite-induced hydroxyl radical
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(·OH) and hydrogen peroxide (H2O2) are the oxidants for Sb(III) oxidation. ·OH is the
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oxidant for Sb(III) oxidation in acidic solutions, and H2O2 becomes the main oxidant
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in neutral and alkaline solutions. ·OH and H2O2 can be generated by the reaction of
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previously existing FeIII(pyrite) and H2O on pyrite in anoxic conditions. The oxygen
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molecule is the crucial factor in continuously producing ·OH and H2O2 for Sb(III)
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oxidation. The efficiency of Sb(III) oxidation was enhanced in surface-oxidized pyrite
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(SOP) suspension, more ·OH formed through Fenton reaction in acidic solutions, but
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Fe(IV) and H2O2 was formed in neutral and alkaline solutions. Under the illumination
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of simulated sunlight, more ·OH and H2O2 were produced in the pyrite suspension,
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and the oxidation efficiency of Sb(III) was remarkably enhanced. In conclusion,
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Sb(III) can be oxidized to Sb(V) in the presence of pyrite, which will greatly
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influence the fate of Sb(III) in the environment.
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INTRODUCTION
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Antimony (Sb) is widely used in a variety of industrial products (approximately
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1.4×105 tons each year), such as flame retardants, batteries, alloys, and catalysts.1 The
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exploitation and utilization of Sb results in elevated concentrations of Sb in many
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soils and waters, especially around mining and smelting areas2,
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ranges.4 It is considered a priority pollutant by the Environmental Protection Agency
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of the United States (USEPA) and the Council of the European Communities.5, 6 In the
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natural environment, Sb(V) is the predominant species and exists as Sb(OH)-6 in oxic
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environments, and Sb(III) primarily occurs as Sb(OH)3 and is more stable and toxic
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under anoxic conditions between pH 2 and 10.1, 7 The mobility and toxicity of Sb
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greatly depend on its oxidation state. Therefore, understanding Sb speciation in soil
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and aquatic systems is important for assessing its fate and risk.
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and at shooting
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Sb(III) can react with relevant substances when it enters into the natural
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environment. Several studies have previously been conducted on the oxidation of
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Sb(III) to Sb(V), including dissolved molecular oxygen and hydrogen peroxide,8, 9
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iodate,10 natural minerals(Fe and Mn oxyhydroxides, etc)11-15 and humic acid.16
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Notably, the natural minerals exist widely and have large surface areas, acting as
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strong oxidants in transforming Sb(III) to Sb(V). Mn oxyhydroxides can directly
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oxidize Sb(III). Wang et al. (2012)13 investigated a synthetic manganite (γ-MnOOH)
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for Sb(III) oxidation in aqueous suspensions, and the oxidation of Sb(III) by
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manganite occurred on a time scale of minutes. Fe oxyhydroxides can catalyze Sb(III)
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oxidation by dissolved oxygen14 or light11. Photo-induced oxidation of Sb(III) on 4
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goethite was investigated recently, the results indicated that significant amounts of
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Sb(III) were oxidized to Sb(V) when the suspension was exposed to light.11 Hence,
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oxidation on mineral surfaces is an important reaction in the environmental fate of
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Sb(III).
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Pyrite (FeS2) is the most abundant natural sulfur mineral in the earth's crust. A
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series of complex reactions occur upon addition of pyrite to O2-free water; many free
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radicals and oxidants, such as the hydroxyl radical (·OH) and hydrogen peroxide
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(H2O2), are formed (Table 1). Numerous studies have reported the degradation of
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various organic pollutants by pyrite-induced ·OH and H2O2 in the environment.17-24
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Simultaneously, pyrite is also associated with stibnite; thus, the speciation of Sb may
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also be influenced by pyrite, which is widespread in soils, waters and sediments.
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Therefore, we deduce that pyrite-induced H2O2 and ·OH could oxidize Sb(III) to
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Sb(V). However, to our knowledge, no studies have examined the reaction of Sb(III)
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with pyrite or other sulfide minerals.
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The present study aimed to investigate the reaction between Sb(III) and pyrite and
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further study the reaction mechanisms under specific conditions (oxic, anoxic,
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surface-oxidized pyrite (SOP) and simulated solar light). Understanding the
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mechanism of Sb(III) oxidation by pyrite is helpful in clarifying the fate and
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geochemical cycling of Sb in the environment.
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MATERIALS AND METHODS
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Reagents and Materials 5
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Antimony trioxide, Sb2O3 (99.999%) and pyrite (naturally occurring mineral) were
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obtained from Alfa Aesar (Johnson Matthey, USA). HCl, KOH, H2O2, FeSO4,
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Fe2(SO4)3, tert-Butyl alcohol (TBA), thiourea, ascorbic acid and KBH4 are guaranteed
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reagent (GR) grade, 2-Morpholinoethanesulfonic Acid (MES), K2B4O7, H2BO3, citric
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acid monohydrate, phenanthroline, CH3COONH4 and CH3COOH are analytical
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reagent
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Chemical Reagent Co. Ltd. (Shanghai, China). Catalase from bovine liver was
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purchased from Sigma-Aldrich (St. Louis, MO, USA). Sb(III) stock solution was
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prepared by dissolving Sb2O3 in 2 mol L-1 HCl. All aqueous solutions were prepared
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using deionized water treated with a Milli-Q water purification device (Millipore),
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stored at 4°C and protected from light.
(AR) grade. The above chemicals were purchased from Sinopharm
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Pyrite was ground and sieved to a 200 mesh powder, washed with 1 M HCl to
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remove surface oxidation layers, rinsed three times with deoxygenated deionized
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water and dehydrated with ethanol, and dried and stored in a closed vial under a N2
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atmosphere. Then, the pyrite was exposed to air at room temperature for 2 weeks, and
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SOP was obtained. BET surface areas of the pyrite and SOP were 0.214 ± 0.021 and
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0.281 ± 0.025 m2 g-1, respectively. The pyrite and SOP were characterized by scanning
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electron microscopy (SEM), energy dispersive spectrometer (EDS) (S-4800, Hitachi,
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Japan) and X-ray diffraction (XRD) (X' Pert PRO MPD, PANalytical, Holland)
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analysis (Figures S1). X-ray photoelectron spectroscopy (XPS) (ESCALAB 250Xi,
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Thermofisher, Japan) was performed for evidence of the speciation of Sb (Figures S4)
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and Fe (Figures S6) on the surface of pyrite and SOP, respectively. The XPS spectra 6
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for Sb 3d5/2 and Fe 2p3/2 were fitted using XPSPEAK41 software.
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Oxidation Reaction Experiment in the Dark
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Sb(III) oxidation in pyrite suspension under oxic and anoxic conditions were
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investigated at pH 3, 5, 7 and 9 under oxic and anoxic conditions with 20 µM Sb(III)
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and 0.25 g L-1 pyrite. In anoxic experiments, solutions were purged with N2 (99.9%)
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for at least 6 h to ensure that oxygen was excluded. The quenching experiments were
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performed with 200 mM TBA or 0.075 g L-1 catalase, 20 mM Sb(III) and 0.25 g L-1
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pyrite to investigate pyrite-induced ·OH and H2O2 on Sb(III) oxidation at pH 3, 5, 7
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and 9. In addition, 20 µM Sb(III) oxidation experiments by 50 µM H2O2, 20 µM Fe(II)
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or 20 µM Fe(III) were performed at pH 3, 5, 7 and 9 in the absence of pyrite,
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respectively.
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All solutions were prepared in a background of 12 mM KCl. Acidic solution (pH 3)
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was prepared via small additions of 0.1 M HCl or KOH, circumneutral solution (pH 5
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and 7) and alkaline solution (pH 9) were kept constant using 2 mM MES/KOH and
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0.5 mM K2B4O7/H2BO3 buffer solution, respectively. The experiments with 200 mL
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solution were performed in triplicate in 250 mL polyethylene bottles and shaken at
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180 rpm at 25°C for 6 h, the suspensions were taken out at selected time intervals. For
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the analysis of dissolved Sb(III) and total Sb (TSb) in the presence of pyrite, 5 mL
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aliquots were removed with a plastic syringe and filtered through a 0.22 µm cellulose
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acetate membrane immediately and then diluted into polyethylene bottles. For the
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analysis of Sb(III) and TSb in H2O2, Fe2+ and Fe3+ experiments, 1 mL aliquots were
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removed to 1ml 1:1 HCl to dissolve solid phase and then diluted into polyethylene 7
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bottles. The samples was immediately analyzed.
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Photo-oxidation Reaction Experiment
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A diagram of the reactor is shown in Figures S2. Batch photo-induced oxidation
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experiments were conducted in 500 mL beakers. During each experiment, 200 mL of
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solution was placed in a beaker containing given concentrations of pyrite and Sb(III).
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The initial Sb(III) concentration was 20 µmol L-1, and the pyrite concentration was
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0.25 g L-1. The control experiment in the dark was conducted by aluminum foil
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coverage. A 500 W long-arc xenon lamp (35 cm×1 cm) (Shanghai Jiguang Special
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Lighting Appliance Factory, China) with a lamp cover used to irradiate visible light
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(λmax = 400~700 nm) was placed above the beaker. The reaction temperature was
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maintained at 25 ± 2°C by cold water circulation. Then, the suspension was exposed
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to light for 2 h. Samples were taken at selected time intervals and filtered through
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0.22 µm cellulose acetate membranes for analysis. All of the experiments were
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conducted in triplicate.
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Analysis
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TSb and Sb(III) were selectively analyzed using a hydride generation-atomic
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fluorescence spectrometer (HG-AFS) (AFS 9700 Titan Instrument Co. Ltd., Beijing,
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China); 10% (m/V) thiourea and ascorbic acid were used to reduce Sb(V) to Sb(III).
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The conditions for the hydride generation system were as follows: carrier solution 5%
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(V/V) HCl, 2% (m/V) potassium borohydride and 0.5% (m/V) potassium hydroxide.
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The concentration of Sb(III) was determined after adding 10% citric acid for masking
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the Sb(V) signal with the same hydride generation conditions as the determination of 8
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TSb. Sb(V) concentration was calculated as the difference between the concentration
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of TSb and Sb(III). The calibration curves were obtained from 0 to 50 µg L-1 of TSb.
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The relative standard deviation was 2%, obtained in 10 measurements performed per
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day. The method detection limit was 0.01 µg L-1. Aliquot samples were diluted in 6
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mol L-1 hydrochloric acid to fit within the working range.
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For Fe(II) analysis, 2 mL of solution was taken and added to 10 mL centrifuge
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tubes. Then, 0.2 mL of 2 g L−1 phenanthroline and 1 mL of CH3COONH4-CH3COOH
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buffer solution were added to the centrifuge tubes. After dilution to 10 mL and 15 min
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of color development, the absorbance of the samples at 510 nm was analyzed on a
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spectrophotometer (Shimadzu UV-2500) using a 5 cm quartz cuvette.25, 26
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Electron spin resonance (ESR) measurements
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To verify the existence of hydroxyl radicals in pyrite solutions, ESR measurements
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were performed on a Bruker EleXsys E500 EPR spectrometer (Germany) at room
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temperature (25 ± 1°C) using 50 µL capillary tubes. Typical instrument settings were
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as follows: sweep width 100.0 G, power attenuation 13.0 dB, modulation amplitude of
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2 G and sweep time 40 s. The ESR signals were recorded from the samples for 320 s
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at pH 3, 6 and 9 in the dark, and a 500 W high-pressure mercury lamp was used to
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investigate the existence of hydroxyl radicals in the presence of light.
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RESULTS AND DISCUSSION
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Sb(III) oxidation by pyrite under oxic and dark conditions
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To investigate Sb(III) oxidation in pyrite solutions, the kinetics experiments were 9
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conducted at pH 3, 5, 7 and 9 (Figure 1). The concentration of Sb(V) increased with
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increasing reaction time (Figure 1a), indicating that Sb(III) was oxidized to Sb(V) in
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the presence of pyrite. Figure 1b further exhibits the Sb(III)/TSb ratio in pyrite
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solution. The Sb(III)/TSb ratio decreased with increasing reaction time, and the
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oxidation efficiencies of Sb(III) were gradually enhanced as the solution pH ranged
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from acid to alkaline. However, in the absence of pyrite, no significant Sb(III)
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oxidation was observed in 6 h reactions (Figure S3).
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The concentrations of TSb in solution decreased with increasing reaction time
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(Figure 1a), which illustrated that a portion of the Sb was adsorbed by pyrite; the
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adsorption amount increased gradually as pH ranged from acidic to alkaline. To verify
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the species of Sb adsorbed on pyrite, XPS analysis was performed at pH 5 after a 6 h
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reaction. The Sb 3d5/2 spectrum was collected and analyzed to determine the state of
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Sb on pyrite (Figure S4). Two peaks of 531.0 and 532.3 eV binding energies for Sb(III)
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and Sb(V), respectively, were obtained from the original Sb 3d5/2. The majority of Sb
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adsorbed on the pyrite surface was Sb(V); the quantitative analysis showed that 73.0%
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of the Sb on pyrite existed as Sb(V). The experimental results confirmed that Sb(III)
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was oxidized to Sb(V) in both solution and on the surface of pyrite.
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Mechanisms on the formation of H2O2 and ·OH by pyrite in oxic aqueous solution
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were investigated systematically in previous studies.27, 28 Sulfur-deficient defects are
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common on pyrite.29, 30 In aerobic conditions, FeII(pyrite) can react with oxygen to form
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FeIII(pyrite) and H2O2 at a sulfur-deficient defect site (Equ. 8, 9); then, FeIII(pyrite) reacts
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with the adsorbed H2O at a sulfur-deficient defect site on the pyrite surface to form 10
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·OH (Equ. 10), and two ·OH combine to produce H2O2 (Equ. 25). Therefore, we
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believe that the pyrite-induced ·OH and H2O2 could oxidize Sb(III).
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The ESR confirmed the existence of ·OH in pyrite suspension under aerobic
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conditions, and the concentrations of ·OH, which were highest at pH 3 and lowest at
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pH 9, decreased with increasing pH (Figure S5). Fe(II) can be released to the solution
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in the process of pyrite oxidation (Equ. 5, 6). Approximately 2.5 µM Fe2+ was found
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in pyrite solution at pH 3 after a 6 h reaction; thus, more ·OH could form via a Fenton
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reaction in acidic solution, but ·OH formation was inhibited as pH increased (Equ. 15,
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16). There was a discrepancy between •OH concentrations and Sb(III) oxidation at
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different pH values. Hence, ·OH is likely not the only oxidant for Sb(III) oxidation,
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and other oxidants must exist. Once ·OH formed, it could rapidly combine to form
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H2O2 (Equ. 25). To investigate the effect of H2O2 on Sb(III) oxidation, Sb(III)
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oxidation experiments by H2O2 were conducted (Figure S6). At pH 3 and pH 5, H2O2
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could not react with Sb(III), but the oxidation of Sb(III) occurred at pH 7 and 9, and
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the oxidation efficiency of Sb(III) was higher at pH 9 than pH 7. Previous studies
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showed that H2O2 was an important and effective oxidant in the process of Sb(III)
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oxidation in alkaline aqueous environment with half-lives of 117 and 11 days at pH 8
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and pH 9 with H2O2 concentrations 10-6 M, respectively8, 9. According to our
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experiments above, we speculated that ·OH and H2O2 could be the main oxidants for
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Sb(III) oxidation in acid solutions and alkaline solutions, respectively, in pyrite
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suspensions.
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To further verify the mechanisms of Sb(III) oxidation by pyrite-induced ·OH and 11
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H2O2, quenching experiments were conducted at pH 3, 5 and 9 by adding tert-butyl
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alcohol (TBA) and catalase into the pyrite suspension (Figure 2). TBA is a widely
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used scavenger of ·OH26,
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peroxide to water and molecular oxygen but also efficiently scavenge ·OH,32 with the
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constant rate of 2.6×1011 M-1 s-1. At pH 3, Sb(III) oxidation was completely quenched
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by TBA. At pH 5, the contribution of ·OH in Sb(III) oxidation was approximately
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73.1%, but at pH 9, this proportion dropped to approximately 5.9%. These results are
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consistent with the ESR results, which implies that ·OH is the main oxidant for Sb(III)
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at low pH. The addition of catalase conspicuously decreased the oxidation efficiency
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of Sb(III) at pH 5 and 9, and the results implied that H2O2 is the key oxidant in Sb(III)
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oxidation in pyrite suspensions in alkaline conditions.
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Sb(III) oxidation by pyrite under anoxic and dark conditions
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. Catalase can not only efficiently convert hydrogen
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The experiments were conducted in anoxic conditions to investigate the effect of
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oxygen on Sb(III) oxidation in pyrite suspension. The oxidation efficiencies of Sb(III)
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under two conditions (oxic or anoxic) at pH 5 are shown in Figure 3. The oxidation
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efficiency of Sb(III) was 37.9% in oxic conditions, whereas it dropped to 15.6% at pH
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5 in anoxic conditions. The results indicated that the presence of oxygen enhances the
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oxidation efficiency of Sb(III) in pyrite suspension. In anoxic conditions, the observed
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rate constant of Sb(III) (Kobs = 0.0010 min-1) was slightly lower than that in oxic
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conditions (Kobs = 0.0012 min-1) in the first 2 h, but an obvious inhibitory effect was
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observed after 2 h (Kobs = 0.0001 min-1).
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In oxic conditions, a cycle between FeII(pyrite) and FeIII(pyrite) exists (Equ. 8-10). 12
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However, in anoxic conditions, the oxidation process of FeII(pyrite) to FeIII(pyrite) cuts off,
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and the production of ·OH and H2O2 is inhibited. However, the oxidation of Sb(III)
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was still occurring in anoxic conditions in the first two hours, implying that ·OH and
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H2O2 could be generated by pyrite in the absence of molecular oxygen. The original
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FeIII(pyrite) can still react with H2O to produce ·OH and H2O2, and once the original
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FeIII(pyrite) is exhausted, the oxidation reaction stops. Hence, molecular oxygen is the
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crucial factor in continuously producing ·OH and H2O2.
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Comparison of SOP and pyrite on Sb(III) oxidation
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Pyrite oxidation is an important process in the geochemical cycles of sulfur and
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iron. In aerobic environments, pyrite is thermodynamically unstable and can be
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oxidized by dissolved molecular oxygen, hydrogen peroxide, or dissolved ferric iron
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(Equ. 5-7). Oxidation of pyrite surfaces may occur upon exposure to atmospheric
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O2.33, 34 Our study also showed that the ferric and ferrous (hydr)oxides were formed
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on the surface of SOP, which were verified by XPS analysis of Fe 2p spectra and EDS
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analysis of SOP (More details were exhibited in Supporting information, Figure S1,
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S7). To compare the difference and mechanism of Sb(III) oxidation by pyrite and SOP,
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experiments were performed in solutions of 0.25 g L-1 pyrite and SOP containing 20
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µM Sb(III) (Figure 4). The oxidation efficiencies of Sb(III) by SOP were enhanced
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compared with pyrite at pH 3, 5, 7, and 9.
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More Fe(II) is released to solution in SOP suspension. Approximately 19 µM Fe(II)
256
was found in the SOP solution at pH 3, whereas nearly no Fe(II) existed at pH 9. To
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investigate the effect of Fe(II) on Sb(III) oxidation, experiments of 20 µM Sb(III) in 13
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the presence of 20 µM Fe2+ were carried out at pH 3, 5, 7 and 9 under oxic and dark
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conditions (Figure S8). At pH 3 and 5, the oxidation efficiencies of Sb(III) were not
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obvious, but at pH 7 and 9, rapid oxidation of Sb(III) occurred, and the oxidation
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efficiencies of Sb(III) increased with increasing pH. The rapid oxidation of Sb(III)
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was due to the formation of H2O2 and Fe(IV) in the presence of Fe(II). Fe(II) was not
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stable in alkaline condition, it was rapidly oxidized by dissolved oxygen, leading to
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the formation of H2O2 (Equ. 11, 12). In addition, Fe(II) and H2O2 can form an
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intermediate (INT), and the INT could lead to formation of Fe(IV) by the outer-sphere
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electron-transfer reaction in neutral and alkaline solutions (Equ. 17-21). The results
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further verified the effect of Fe(II) on Sb(III) oxidation. However, Sb(III) oxidation in
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SOP suspension was also enhanced in acidic conditions, which was attributed to the
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formation of ·OH through a Fenton reaction.
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Notably, SOP has higher surface area, which is about 1.3 times higher than the
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surface area of pyrite, this may be due to the large surface area of formed ferrous and
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ferric (hydr)oxides on the surface of SOP.35 In order to meaningfully compare the
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oxidation reactivity of these two forms of pyrite, Sb(III) oxidation by pyrite and SOP
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was normalized by the surface area (Table 2). SOP showed higher oxidation activity
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for Sb(III), the pyrite oxidation process leads to the transformation of inactive
276
Fe(II)-S to more active Fe(II)-O compounds on the surface of SOP.
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The experiments of Fe(III) on Sb(III) oxidation was also investigated at pH 3, 5, 7,
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and 9. Little Sb(V) was observed in a 6 h reaction in the presence of 20 µM Fe(III)
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and 20 µM Sb(III) (Figure S9); thus, the influence of Fe(III) on Sb(III) oxidation was 14
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negligible in this system.
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Sb(III) oxidation by pyrite in the presence of light
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Pyrite is a semiconductor with a small band gap of 0.95 eV, and it can be activated
283
by light.36,
37
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sunlight on Sb(III) oxidation in pyrite suspension at pH 6 (Figure 5). A small amount
285
of Sb(III) oxidation was observed in the absence of pyrite under light. Remarkably,
286
Sb(III) was completely transformed to Sb(V) after 2 h of reaction in the presence of
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pyrite under light. However, only 35.5% of Sb(III) was oxidized after a 2 h reaction in
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dark. The oxidation efficiency of Sb(III) under light (Kobs = 0.0142 min-1) was
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4.3-fold the rate under dark (Kobs = 0.0033 min-1). The efficiency of Sb(III) oxidation
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by pyrite was greatly increased with exposure to simulated sunlight.
Experiments were performed to investigate the effect of simulated
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When pyrite is illuminated, a photon can be captured and promote an electron from
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a non-bonding iron d-level to the conduction band of pyrite, and hole-electron pairs
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(e-, h+) in valence band are excited on the surface of pyrite. A hole (h+) migrates to the
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pyrite surface, and adsorbed H2O dissociates to form ·OH in pyrite suspension (Equ.
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22, 23).38 This process was confirmed by ESR (Figure S10), which showed an
296
increasing concentration of ·OH under the illumination of light in pyrite suspension.
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Hence, the formation of ·OH via the photochemical process can contribute to the
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promotion of the oxidation efficiency of Sb(III).
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The oxidation efficiency of Sb(III) was altered less by the addition of TBA under
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light than in the dark, indicating that •OH is not the only oxidant generated by light.
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·OH could combine to form H2O2 (Equ. 25). A previous study demonstrated 15
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increasing H2O2 production under the illumination of light in pyrite suspension.39
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Proposed reaction mechanisms
304 305
Based on the above results, possible mechanisms of Sb(III) oxidation in the presence of pyrite were proposed as follows (Figure 6):
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❶ ≡FeII induced ·OH and H2O2 on pyrite surface. ≡FeII can be oxidized by O2 to
307
generate H2O2 and simultaneously provide ≡FeIII at a sulfur-deficient defect site, and
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then, the adsorbed H2O reacts with ≡FeIII, generating adsorbed ·OH and H2O2.
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❷ FeII induced ·OH, H2O2 and Fe(IV) in pyrite solution. The dissolved FeII can
310
react with O2 to generate H2O2. In neutral and alkaline solutions, Fe(IV) can form
311
through the reaction of FeIIOH+ and H2O2. However, in acidic solutions, more ·OH
312
forms through the Fenton reaction.
313 314 315 316
➌ Light-induced ·OH and H2O2. More ·OH and H2O2 form in the illumination of light in pyrite suspension. The pyrite-induced ·OH, H2O2 and Fe(IV) can oxidize Sb(III) to Sb(V) (Equ.1-3).
Environmental implications
317
When redox speciation determinations of Sb are performed in aquatic systems, the
318
results have commonly shown that Sb(V) predominates and that Sb(III) is only found
319
at low concentrations in oxic conditions.1 Significant concentrations of Sb(V) are
320
found in anoxic waters,40 which contrasts with thermodynamic equilibrium
321
predictions. In oxic environments, Sb(V) may come from the catalytic oxidation of
322
Sb(III) by natural Fe oxyhydroxides11, 14 and humic acid16 in the presence of dissolved
323
oxygen or light, or from direct oxidation of Sb(III) by Mn oxyhydroxides13. Whereas 16
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this study indicated that Sb(III) can also be catalytically oxidized by pyrite induced
325
hydroxyl radicals and hydrogen peroxide. The process can greatly increase the
326
mobility and reduce the toxicity of Sb. The oxidation efficiency of Sb(III) is
327
conspicuously enhanced under visible light, and it is 4.3-fold faster than in the dark.
328
This oxidation process may be more important in surface soils or waters. In anoxic
329
environments, pyrite-induced oxidation of Sb(III) can also occur, which could partly
330
explain the existence of Sb(V) in anoxic environments. Therefore, pyrite can greatly
331
influence the fate of Sb(III) and play an important role in the biogeochemical cycle of
332
antimony.
333
ACKNOWLEDGEMENTS
334
This work was supported by the National Science Foundation for Innovative Research
335
Group (Grant No. 51421065), the National Natural Science Foundation of China (21177011)
336
and the Major Science and Technology Program for Water Pollution Control and Treatment
337
(2012ZX07202002). We are grateful to the editors and the anonymous reviewers for their
338
numerous valuable comments and suggestions on our paper.
339
Supporting Information Available
340
Details of SEM, EDS and XRD analysis of pyrite and SOP, photochemical reactor,
341
Sb(III)/TSb changes in solution in the absence of pyrite, XPS analysis of Sb speciation, ESR
342
spectra in dark and light generated by pyrite, Sb(III) oxidation by H2O2, Fe2+ and Fe3+ and
343
XPS analysis of Fe 2p on the surface of SOP. The Supporting information is available free of 17
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charge via the Internet at http://pubs.acs.org/.
345 346 347 348 349 350 351 352 353
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TABLE CAPTIONS
472
Table 1. Chemical Reactions and Corresponding Rate Constants in Pyrite Suspension
473
Table 2. The initial and normalized oxidation rate of Sb(III) by pyrite and SOP
474
FIGURES CAPTIONS
475
Fig. 1 Concentration of TSb and Sb(V) in solution (a) and the oxidation efficiency of
476
Sb(III)/TSb in solution by pyrite (b) under oxic and dark conditions at pH 3, 5, 7 and 9.
477
Inset: pseudo-first-order model for the determination of the observed rate constant Kobs.
478
[FeS2]=0.25 g L-1, [Sb(III)]=20 µM.
479
Fig. 2 Quenching effect of TBA and catalase on Sb(III) oxidation in pyrite suspension under
480
oxic and dark conditions at pH 3 (a), 5 (b) and 9 (c). [FeS2]=0.25 g L-1, [Sb(III)]=20 µM,
481
[TBA]=200 mM, [catalase]=0.075 g L-1.
482
Fig. 3 Sb(III)/TSb ratios under anoxic and oxic conditions in dark at pH 5. Inset:
483
pseudo-first-order model for the determination of the observed rate constant Kobs.
484
[FeS2]=0.25 g L-1, [Sb(III)]=20 µM.
485 486
Fig. 4 Comparison of Sb(III)/TSb ratios in solution between pyrite and SOP at pH 3, 5, 7 and
9 under oxic and dark conditions. [FeS2]=0.25 g L-1, [SOP]=0.25 g L-1, [Sb(III)]=20 µM.
487
Fig. 5 Sb(III)/TSb ratios during oxidation under dark and light conditions in oxic solutions at
488
pH 6. Inset: pseudo-first-order model for the determination of the observed rate constant
489
Kobs. [FeS2]=0.25 g L-1, [Sb(III)]=20 µM, [TBA]=200 mM.
490
Fig. 6 Possible mechanisms of Sb(III) oxidation in pyrite suspension ❶ FeIII(pyrite)-induced
491
oxidation of Sb(III). ❷ FeII(aq)-induced oxidation of Sb(III). ➌ Light-induced ·OH and H2O2.
492
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TABLE 1. Chemical Reactions and Corresponding Rate Constants in Pyrite Suspension Rate coefficients
Equ.
(s-1 or M-1 s-1 M-1 s-1)
ref.
Antimony Oxidations Sb(III)+·OH →Sb(IV)
2
Sb(III) + H2 O2 → Sb(IV)
3
Sb(III) +Fe(IV)→ Sb(IV)+Fe3+
9 41
●
4
41
8 × 109
1
Sb(IV) + O2 → Sb(V) + O2
1.1 × 109
41
Pyrite Oxidations 2+
4SO24
38
+ 4H+
5
2FeS2 + 7O2 + 2H2 O → 2Fe
6
+ FeS2 + 14Fe3+ + 8H2 O →15Fe2+ + 2SO24 + 16H
38
7
+ 2FeS2 + 15H2 O2 → 2Fe3+ + 4SO24 + 2H +14H2 O
42
8
FeII(pyrite) FeII(pyrite) FeIII (pyrite)
+ O2 →
FeII(aq) FeII(aq) FeII(aq) FeII(aq)
O2 → O2 ● O2 + 2H → FeIII aq + H2 O2 ● + III HO2 + H → Feaq + H2 O2 ·OH → FeIII aq + OH
+
Reactions on Pyrite Sulfur-Deficient Defect Sites
9 10
FeIII pyrite
●
+
O2
+
+ 2H →
27
●
+ O2
III Fepyrite
27
+ H2 O2
28
+ H2 O → FeII(pyrite) + ·OH FeII Oxidations
11 12 13 14
+ + + +
43
●
FeIII aq + +
44
1 × 107 1.2 × 106
44
8
45
3.2 × 10 Fenton Reactions
2+
3+
15
Fe
16
Fe OH + H2 O2 → FeIII OH2+ + ·OH + OH
+ H2 O2 → Fe
+ ·OH + OH
+
II
Modified Fenton Reactions 2+
46
17
Fe
18
FeII OH+ + H2 O2 → INT-OH
46
19
INT-OH + H+ → INT
46
20
INT → FeIII OH2+ + ·OH
46
21
INT-OH → Fe(IV)
+ H2 O2 → INT
46
Reactions on Pyrite under Illumination of Light →e
+
22
FeS2 + hv
23
H2 O + ℎ → H+ + ·OH
24
O2 + e- → O2
25
·OH + ·OH → H2 O2
+h
●
Reactive Oxygen Radical Reactions
26
●
HO2 + O2 + H+ → H2 O2 + O2 ●
27
●
●
HO2 + HO2 → H2 O2 + O2
5.2 × 109
47
9.7 × 107
44
8.3 × 105
44
9
6.6 × 10
47
3 × 107
47
●
28
HO2 + ·OH → O2 + H2 O ●
29
H2 O2 + ·OH → HO2 + H2 O
30
O2 + H+ ↔ HO2
31
·OH + TBA → O2 + products
●
1 × 1012/1.58 × 107
●
44
(pKa 4.8) ●-
(3.8~7.6)×108
493 25
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Table 2. The initial and normalized oxidation rate of Sb(III) for pseudo-first-order model by pyrite and SOP Pyrite
SOP
Initial rate Normalized rate
Initial rate Normalized rate
(min-1)
(min-1 m-2)
(min-1)
(min-1 m-2)
pH 3
0.0003
0.0016
0.0007
0.0025
pH 5
0.0012
0.0057
0.0048
0.0171
pH 7
0.0015
0.0068
0.0050
0.0178
pH 9
0.0254
0.1186
0.0575
0.2046
497 498
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Figure 1. Concentration of TSb and Sb(V) in solution (a) and the oxidation efficiency of Sb(III)/TSb in solution by pyrite (b) under oxic and dark conditions at pH 3, 5, 7 and 9. Inset: pseudo-first-order model for the determination of the observed rate constant Kobs. [FeS2]=0.25 g L-1, [Sb(III)]=20 μM.
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Figure 2. Quenching effect of TBA and catalase on Sb(III) oxidation in pyrite suspension under oxic and dark conditions at pH 3 (a), 5 (b) and 9 (c). [FeS2]=0.25 g L-1, [Sb(III)]=20 μM, [TBA]=200 mM, [catalase]=0.075 g L-1.
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Figure 3. Sb(III)/TSb ratios under anoxic and oxic conditions in dark at pH 5. Inset: pseudofirst-order model for the determination of the observed rate constant K obs. [FeS2]=0.25 g L-1, [Sb(III)]=20 μM.
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Figure 4. Comparison of Sb(III)/TSb ratios in solution between pyrite and SOP at pH 3, 5, 7 and 9 under oxic and dark conditions. [FeS2]=0.25 g L-1, [SOP]=0.25 g L-1, [Sb(III)]=20 μM.
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Figure 5. Sb(III)/TSb ratios during oxidation under dark and light conditions in oxic solutions at pH 6. Inset: pseudo-first-order model for the determination of the observed rate constant Kobs. [FeS2]=0.25 g L-1, [Sb(III)]=20 μM, [TBA]=200 mM.
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Figure 6. Possible mechanisms of Sb(III) oxidation in pyrite suspension ❶ ≡FeIII induced oxidation of Sb(III). ❷ FeII induced oxidation of Sb(III). ➌ Light-induced ∙OH and H2O2.
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